Phase Peroxy Radical Reactions

Laboratory Studies of Atmospherically Important Gas- Phase Peroxy Radical Reactions

Thesis by Lance E. Christensen

In partial fulfillment of the requirements for the Degree of Doctor of Philosophy

California Institute of Technology

© 2002 Lance Christensen All Rights Reserved

Abstract

Peroxy radicals (HO2, RO2) are important intermediates in Earth’s atmosphere.

They are intermediates in the oxidation of alkanes and CO in combustion and atmospheric chemical processes. In earth’s atmosphere, the rates of their self and cross reactions are often the dominant loss processes when NOx concentrations fall below tens of pptv. These reactions have proven difficult to study in laboratory experiments, due to complex secondary chemistry and ambiguities in radical detection.

This thesis describes a new laser-photolysis apparatus to measure the rates of peroxy radical reactions under atmospheric conditions that employs simultaneous UV direct absorption and IR wavelength-modulation spectroscopy to detect the peroxy radicals. Prior kinetic measurements of gas-phase peroxy radical reactions have typically employed flash-photolysis methods coupled with detection of the radicals only by UV absorption spectroscopy. However, uncertainties can arise because several different species often contribute to the absorption signal. The IR channel provides an independent means of monitoring HO2 radicals by detection of specific rovibrational transitions.

With this apparatus, the rates of the reactions HO2 + NO2, HO2 + CH3O2, CH3O2

+ CH3O2, and HO2 + HO2 were studied at temperatures from 219 K to 300 K. Our measurements have, in some cases, led to significant revision of previously accepted rate constants, mechanisms, or product yields, especially at conditions relevant to the upper atmosphere. The new rate coefficients for the HO2 + HO2 reaction are shown to account

ii

for a long-standing discrepancy in modeled vs. observed hydrogen peroxide in the stratosphere.

A key finding has been the observation that many previous measurements of HO2

reactions at low temperatures have suffered from problems due to complexation between HO2 and methanol, a precursor used to generate HO2. Direct kinetic evidence is presented for the formation of the HO2·CH3OH complex; the rate coefficients, equilibrium constant, and enthalpy of reaction for HO2 + CH3OH, ↔ HO2·CH3OH were measured. These results are the first direct study of the chaperone effect proposed to explain the enhancement of the observed rates of the HO2 self-reaction by hydrogen- bonding species.

The effects of methanol enhancement on the HO2 + NO2, HO2 + CH3O2, CH3O2 + CH3O2, and HO2 + HO2 reaction rates were measured. For the HO2 + NO2 reaction, overlapping, time-dependent signals in the UV due to the equilibrium between NO2 and N2O4 were observed that may not have been properly accounted for in previous measurements. Other studies of NO2 reactions conducted at temperatures below 250 K may be subject to similar errors. In the CH3O2 + CH3O2 reaction, detection of HO2

products has raised questions concerning the product yields and reaction mechanisms.

Table of contents

Chapter Page

Chapter 1: Measurements of the Rate Constant for HO2 + NO2 + N2 → HO2NO2 + N2 Using Infrared Wavelength- Modulation Spectroscopy and UV-Visible Absorption

Spectroscopy 1

Chapter 2: Kinetics of HO2 + HO2 → H2O2 + O2:

Implications for Stratospheric H2O2 45

Chapter 3: The Methanol Chaperone Effect on HO2

Reactions 62

Chapter 4: Kinetics of CH3O2 Reactions 93

Chapter 5: Experimental Details 117

iv

List of Tables

Chapter 1. Measurements of the Rate Constant for HO2 + NO2 + N2 → HO2NO2 Using Infrared Wavelength-Modulation Spectroscopy and UV- Visible Absorption

Spectroscopy.

1.1 Cross sections for various species... 28

1.2 Relevant reactions... 29

1.3 Fitted values at different temperatures... 30

1.4 Fitted values for Troe equation... 31

Chapter 2. Kinetics of HO2 + HO2 → H2O2 + O2: Implications for Stratospheric H2O2. 2.1 Experimental conditions... 57

Chapter 3. The Methanol Chaperone Effect of HO2 Reactions. 3.1 Relevant reactions... 82

3.2 Experimental conditions... 83

3.3 Values of Keq, k7, and k-7... 84

Chapter 4. Kinetics of CH3O2 Reactions. 4.1 Experimental conditions... 104

4.2 Reaction mechanism... 105

4.3 Values of k2 and α... 106

4.4 Measurements of k1... 107

List of Figures

Chapter 1. Measurements of the Rate Constant for HO2 + NO2 + N2 → HO2NO2 Using Infrared Wavelength-Modulation Spectroscopy and UV- Visible Absorption

Spectroscopy.

1.1 Experimental apparatus... 32

1.2 Simulated UV absorbances at 369.50 nm using FACSIMILE... 33

1.3 Decay of [HO2] due to the HO2 + NO2 reaction at different [CH3OH] at 231 K... 34

1.4 k′ versus [CH3OH] for various [NO2] at 231 K, 100 Torr... 35

1.5 k′o versus [NO2]...36

1.6 k1 versus T compared with the NASA recommendation and expected rate if the HO2 + NO2 reaction were studied using [CH3OH] = 3 × 10¹⁵ molecules cm^-3... 37

1.7 k″ versus [NO2]... 38

1.8 k″o and k^† versus T^-1... 39

1.9.1 Comparison of UV and IR signals at 298 K... 40

1.9.2 Comparison of UV and IR signals at 231 K... 41

1.10 k′ versus [NO2]... 42

1.11 Measured rates of k1 from the present work using [CH3OH] = 4 × 10¹⁴ molecules cm^-3 compared with the NASA recommended values... 43

1.12 Comparison of NASA recommended k1 versus k1 from new parameterization employing the kinetic data from this work with previous studies in which only measurements in which the influence of methanol was insignificant were used... 44

Chapter 2. Kinetics of HO2 + HO2 → H2O2 + O2: Implications for Stratospheric H2O2. 2.1 Plot of kobs as a function of [CH3OH] at 231 K and 295 K... 58

2.2 Plot of the rate constant of reaction (1) as a function of inverse temperature at 100 Torr... 59

2.3 Plot of k″ as a function of inverse temperature from the present study at 100 Torr and from the Andersson et al. study at 760 Torr... 60

2.4 Measured and modeled profiles of H2O2 VMR for two seasons near Ft. Sumner, NM (34.5°N)... 61

vi

Chapter 3. The Methanol Chaperone Effect of HO2 Reactions.

3.1 Time dependence of HO2 signal at different methanol

concentrations at 251 K, 100 Torr... 85

3.2 The Dependence of [HO2]o/[HO2]eq on Methanol Concentration at 231 K and 261 K... 86

3.3 Comparisons of experimentally measured and theoretically calculated Kc... 87

3.4 Decay of [HO2] from the reaction HO2 + CH3OH →^M HO2·CH3OH at 240 K, 100 Torr... 88

3.5 Comparisons of the observed rate coefficient for the HO2 self-reaction between the IR and UV detection channels at two different temperatures... 89

3.6.1 Typical example of the IR signal at 231 K, 100 Torr... 90

3.6.2 Typical example of the UV signal at 231 K, 100 Torr... 91

3.7 kobs,ir and kobs,uv versus methanol concentration at 231 K, 100 Torr... 92

Chapter 4. Kinetics of CH3O2 Reactions. 4.1.1 Time dependence of the [HO2] at different [H2]/[CH4] at 231 K, 100 Torr... 108

4.1.2 Time dependence of [CH3O2] at different [H2]/[CH4] at 231 K, 100 Torr... 109

4.2 Natural log plots of data acquired at 252 K, 100 Torr in which the time dependence at [H2]/[CH4]=0 have been subtracted... 110

4.3 Arrhenius Plot of k2 Versus T^-1... 111

4.4.1 Fits using FACSIMILE to the time dependences of [HO2] and [CH3O2] at 296 K... 112

4.4.2 Fits using FACSIMILE to the time dependences of [HO2] and [CH3O2] at 231 K... 113

4.4.3 Comparisons of [HO2] from the CH3O2 self-reaction at 296 K and 231 K... 114

4.5 Possible reaction pathways... 115

4.6 Possible reaction pathway for formation of HO2... 116

Chapter 5. Experimental Details.

5.1 The main reaction cell... 134

5.2 Probe input and aluminum block... 135

5.3 Excimer input and aluminum block... 136

5.4 Joiner for reaction cell and excimer input aluminum block... 137

5.5 Pre-cooling cell... 138

5.6 Photolysis volume... 139

5.7 Calculated HO2 concentration profiles at 100 Torr, 298 K at different times after the photolysis event... 140

5.8 Modeled mass transport rates...141

5.9 Herriott mirrors... 142

5.10 Diode laser beam placement on Herriott mirrors... 143

5.11 Modulation and detection electronics... 144

5.12 HO2 spectrum near 6638.2 cm^-1 as a function of input current to the diode laser... 145

5.13 Two water lines acquired by a DFB diode laser obtained from the Microdevices laboratory at JPL... 146

5.14 Comparisons of HO2 1²A′ and ← X²A″ and O-H overtone transitions... 147

1

Chapter 1: Measurements of the Rate Constant for HO

₂

+ NO

₂

+ N

₂

→ → → → HO

₂

NO

₂

+ N

₂

Using Infrared

Wavelength-Modulation Spectroscopy and UV-Visible Absorption Spectroscopy

1.1 Introduction

The reaction between HO2 and NO2 has been the subject of numerous laboratory studies^1-10 and proceeds as

M

2 2 2 2

HO + NO → HO NO (1)

From the upper troposphere through the middle stratosphere, the thermal lifetime of HO2NO2 is sufficiently long that reaction with OH is a significant loss process for HO2NO2. This establishes a NO2 driven catalytic cycle that is an important sink for HOx.¹¹

M

2 2 2 2

HO + NO → HO NO (1)

2 2 2 2 2

OH + HO NO →H O + NO + O (2)



2 2 2

OH + HO → H O + O (3)

Measurements of HO2NO2 from space¹² and balloon-borne¹³ platforms have enabled researchers to test our understanding of atmospheric processes involving HO2NO2. Accurate measurements of k1 are thus necessary to correctly describe the chemistry of this region of the atmosphere.

The most comprehensive studies of reaction (1) measured the time dependence of [HO2] with UV spectroscopy and utilized CH3OH as a precursor for HO2.^8-10 These studies have had the greatest influence on current recommendations.^14,15 In these previous studies, the rate of reaction (1) was measured under conditions where an appreciable fraction of HO2 would be hydrogen-bonded to CH3OH, namely low temperatures (<

273 K) and/or high [CH3OH]. It has been shown^16,17 that CH3OH can significantly enhance the observed rate of the reaction

bi-molecular

2 2 ter-molecular 2 2 2

HO + HO → H O + O (4)

under these conditions. A similar enhancement for reaction (1) might also be expected.

This would suggest that the NASA recommended rates are too high at low temperatures.

This paper details kinetic studies of reaction (1) using simultaneous UV and IR detection. The effect of CH3OH on reaction (1) was measured. Detection of HO2 in the IR provided a method of measuring k1 that avoided overlapping absorptions from several species, a problem associated with measurements in the UV. In addition, the use of heterodyne detection for the IR channel resulted in considerably improved signal-to-noise

3

1.2 Experimental 1.2.1 Apparatus

Figure 1.1 is a schematic diagram of the experimental apparatus. A XeCl pulsed excimer laser (308 nm) was used to photolyze either F2 or Cl2, which reacted with other reagents to form the species of interest. The concentrations of the species of interest were monitored with simultaneous IR heterodyne and UV-visible direct absorption spectroscopy.

The reaction cell was a 175 cm long, 5 cm diameter Pyrex cylinder supported at each end by aluminum chambers. An insulated jacket surrounded the reaction cell through which flowed methanol chilled by a liquid-nitrogen cooled heat exchanger.

Thermocouples located inside the reaction cell allowed the temperature to be measured to within ± 1 K. Reagent gases were cooled in a meter-long mixing tube prior to entering the main reaction cell. They entered from the middle of the main reaction cell and flowed towards the outlet ports. N2 confinement gas flowing from both aluminum chambers restricted the reactants to a region 135 ± 1 cm long between the outlet ports. Tests were performed to ensure the extent of confinement by flowing gas mixtures containing known amounts of Cl2 and NO2 through the reagent entrance port. In these tests, the Cl2

absorbance at 330 nm and NO2 absorbance at 369.50 nm was measured. The effective path length was calculated using a Beer’s Law analysis and tabulated absorption cross sections.¹⁸ These tests were conducted over the range of pressures and flow rates utilized in the experiment. They confirmed that the reagent gases were contained between the two

exit ports with an effective path length matching the separation between the middle of the two ports to within 1 cm.

The excimer photolysis pulse entered the cell through a CaF2 window on one of the aluminum chambers. The 20 ns pulses had a 2 cm by 1 cm rectangular cross section.

The pulse energy ranged from 60 mJ to 150 mJ. The pulses passed through the middle of the cell, creating a 2 cm by 1 cm by 138 cm photolysis region. An unstable optical resonator configuration was used in the excimer laser to ensure good collimation of the photolysis beam.

Light from a deuterium (D2) lamp and an IR diode laser entered the apparatus through a 30′ wedged CaF2 window on the other aluminum chamber. Light from the D2

lamp made one pass through the photolysis volume and was focused onto the entrance slit of a monochromater (Acton SpectraPro 300i). A PMT was mounted at the exit slit.

Baffles in both aluminum chambers ensured that only UV light that had sampled the photolysis volume entered the monochromater.

For the present experiment, several species were formed which absorb in the UV.

The detectable species were HO2, NO2, N2O4, H2O2, HO2NO2, ClNO2, and ClONO. Their cross sections at various wavelengths are given in Table 1.1. The monochromater was set to 369.50 nm for experiments conducted at temperatures of 230 K and higher and 381.875 nm for experiments at 219 K. The 369.50 nm NO2 cross section at 298 K was 5.23 × 10^-19 cm² with a temperature dependence of -1.1 × 10^-22 cm² K^-1.^19,20 The 381.875 nm NO2 cross section at 298 K was taken to be 5.62 × 10^-19 cm² with a temperature dependence of -8.7 × 10^-23 cm² K^-1.^19,20 The reason for the change in UV

5

The IR source was a 3 mW distributed-feedback (DFB), continuous wave tunable diode laser manufactured by the JPL Microdevices Laboratory. The laser current was modulated at 6.80 MHz through an external bias tee. The beam passed through a small opening in a gold-coated mirror with a 2032 mm radius of curvature located in one aluminum chamber and impinged on a similar mirror in the other chamber positioned 1820 mm from the input mirror. These two mirrors formed a Herriott cell^21,22 that folded the IR beam, resulting in 30 passes through the photolysis volume. The beam was inside the photolysis volume for approximately 1/2 the length of a single pass between mirrors.

The effective path length of the IR beam was approximately 2000 cm as determined by visual inspection of where overlap occurred. This was maximized by placing the Herriott mirrors as close to the path of the excimer beam as possible. The signal from the InGaAs photodiode detector was demodulated at 13.6 MHz (2f detection) and low-pass filtered.

The filter frequency was determined by the timescale of the reaction. Typically, bandwidths greater than a factor of 5 over the pseudo-first-order HO2 loss rate were employed. Minor adjustments of the amplitude of modulation were required to optimize the signal when the pressure and temperature of the cell was varied.

The diode laser emitted light in the region between 6620 cm^-1 and 6645 cm^-1, depending on the injection current and temperature of the diode laser. The lower frequency limit was determined by the maximum temperature the diode laser chip could be held at. For emission at 6620 cm^-1, the temperature had to be set at around 60 ºC. At these temperatures, the lifetime of the laser is drastically reduced. Further, d(Power)/d(Current) becomes appreciably non-linear, introducing a significant amount of noise into the IR detection channel. The upper limit of the laser emission frequency was

to prevent condensation of ambient water when the diode laser was cooled below 5 ºC.

The linewidth of the laser was approximately 20 MHz.²³ This was verified by deconvolving H2O transitions at low pressure (< 200 mTorr). For the present study, an HO2 transition at 6638.2 cm^-1 was probed. This line is assigned to the ^qQ2 transition (a band head) of the first overtone of the O-H stretch.²⁴ Another diode laser that emitted near 7000 cm^-1 (JPL Microdevices Laboratory) was also employed in the experiment but only for a limited number of experiments at room temperature. This diode laser probed transitions to the low-lying electronic state of HO2 (²A′ ← ²A″). No differences in the measured kinetic parameters were observed between the two lasers. For HO2, direct absorption measurements have suggested that the integrated band strength of the overtone transitions absorb are stronger than the electronic transitions.²⁵ The cross-section of the

qQ2 line at 100 Torr, 298 K was estimated to be (5 ± 3) × 10^-20 cm². This was determined by observing that its signal was similar in strength to several of the strongest lines near 6627 cm^-1, which have been assigned to the P-branch of the K″ = 0 stack.²⁴ These lines have been observed to have cross sections between (1 –10) × 10^-20 cm^-2 near 60 Torr.²⁶ The highest concentrations of HO2 employed in the present experiment were around 1 × 10¹⁴ molecules cm^-3. The absorbance for a pathlength of 2000 cm is approximately 0.01.

The difference between Beer’s law analysis and simply correlating the IR signal with [HO2] is less than 1% at the maximum [HO2].

The IR and UV beams have different geometric paths, and consequently probe different regions of the photolysis volume. The IR beam probes the central half of the photolysis volume. The UV beam probes the whole length of the photolysis volume. A

7

simultaneous measurements of the HO2 + HO2 reaction under conditions where the concentrations of species that can hydrogen-bond with HO2 is negligible. The formula

( )

-1

o

( ) - 1 2

S t b a t

S b

 

= + + + ⋅ ⋅  (5)

was employed to study second-order reaction kinetics. S(t) is the signal at either detector as a function of time, So represents the signal extrapolated to time = 0, b represents a constant baseline offset, and a represents the second-order rate constant in units of S(t)^-1 s^-1. For UV measurements, S(t) was in units of absorbance. For IR measurements, S(t) was in units of V. The product a·So for the UV and IR should be equivalent for each experiment and is units of s^-1. The value of a obtained from UV measurements was corrected for the contribution of H2O2 by multiplying its value by ^{2 2}

2

H O HO

1- σ

2⋅σ , following the procedure outlined by Kircher and Sander.²⁷

Simultaneous IR and UV rate measurements of reaction (4) were used to calibrate the IR signal. For rate measurements with the IR probe, calibration of the probe signal was necessary. This was accomplished by simultaneously measuring the time decay of HO2 for HO2 + HO2 with the IR and UV probes, employing F2-photolysis. The path length of the beam was 135 cm. Because the cross section and the path length were known, the UV measurement provided a second-order rate constant in units of cm³ molecule^-1 s^-1. The IR probe measured a second-order rate constant in units of V^-1 s^-1. The ratio of the rate constants gave the scaling factor used to translate the IR signal from

Volts to units of molecules cm^-3. This value ranged between (1 - 6) × 10¹⁷ molecules V^-1 at the RF port of the demodulation mixer.

The photolysis volume was centrally located and wall reactions were not a concern. However, transport of reactive species from the photolysis volume into the surrounding gas by turbulent mixing was an important consideration. The reactions C2H5O2 + C2H5O2 (kEtO2), and HO2 + HO2 were studied at [C2H5O2] < 1×10¹³ molecules cm^-3 and [HO2] < 5 × 10¹¹ molecules cm^-3, respectively. At these concentrations, 2·kEtO2·[C2H5O2] and 2·k4·[HO2] < 2 s^-1, and other loss processes, such as turbulent mixing, can compete with loss due to chemical reaction. The measured rates for these reactions were dependent on the residence time of the precursor gases, indicating that turbulent mixing affected measured kinetics. As the residence time was increased, the measured rates approached the predicted rates asymptotically. The measured first-order loss due to turbulent mixing was between (2 - 5) s^-1 for a residence time of 15 s at 298 K.

Turbulent mixing affects decreased with decreasing temperature. The residence time was adjusted so turbulent mixing had less than a 5% effect on measured rates. The effect of diffusion was found to be negligible compared to turbulent mixing.

Calibrated flows of reagent gases were mixed prior to entering the cell. Flow conditions were adjusted so that the cell residence time was 3-10 s, approximately equal to the interval between photolysis laser pulses. HO2 was formed from the reaction sequence

Cl2→^hv 2 Cl (6)

9

2 2 2 2

CH OH + O → HO + CH O (8)

The concentrations (molecules cm^-3) of the reagents were Cl2: (2 - 6) × 10¹⁵; He: (2 - 5) × 10¹⁶; CH3OH: (1 - 3) × 10¹⁴; O2: (2 - 7) × 10¹⁷; NO2: (6 - 50) × 10¹⁴. The buffer gas was N2 for all experiments. The Cl2 and He came from a mixed cylinder of 10.0% Cl2 (99.5%

purity) in He (99.999%). N2 (99.9993%) was bubbled through CH3OH (A. C. S. Reagent Grade) in a temperature-controlled saturator to obtain the desired [CH3OH]. NO2 was prepared by mixing NO (99% purity) with a large excess of O2 (99.996%) and allowing the mixture to stand for a day. All gases were acquired from Air Products and Chemicals, Inc. except NO, which was acquired from Matheson Tri-Gas, Inc. The maximum [HO2] formed in an individual experiment, denoted [HO2]max, was (5 - 8) × 10¹³ molecules cm^-3.

The reaction between HO2 and NO2 was studied under pseudo-first-order conditions with the value of ^{2 o}

2 max

[NO ]

[HO ] between 60 and 500. The reaction was studied between 40 Torr and 200 Torr and 219 K to 295 K. Contributions to measured NO2

absorbance at 369.50 nm and 381.875 nm from HO2, H2O2, HO2NO2, and Cl2 were less than 2% at all temperatures. At temperatures below 240 K for the concentrations of NO2

employed in the present experiment, a significant fraction of NO2 dimerized to N2O4. At 219 K, ^{2 4}

2

[N O ]

[NO ] reached values as high as 0.8. In order to maintain ^{2 4 2 4}

2 2

σN O [N O ] σNO [NO ]

⋅

⋅ <

0.05, NO2 was monitored at 381.875 nm for experiments at 219 K.

The excimer laser photolyzed a fraction of NO2 to produce NO + O. The fraction was determined to be 0.0028 ± 0.0002. This was measured by photolyzing NO2 in the presence of O2 and observing O3 formation. Davidson et al.²⁰ has shown that at 219 K,

the NO2 cross-section at 308 nm is only 2% higher than at room temperature, indicating that the fraction of NO2 dissociated by photolysis was approximately the same at all temperatures. Knowledge of [NO] produced by photolysis of NO2 was important because the reaction

2 2

HO + NO → OH + NO (9)

can affect the measured rate of decay of HO2. The effect of reaction (9) was ascertained by employing the kinetics modeling program FACSIMILE²⁸ and the NASA recommended¹⁸ values for the rate constants of reactions listed Table 1.2. At 50 Torr and 295 K, the ratio of the observed rate constant to the actual rate constant was calculated to be 1.06. As pressure increases and temperature decreases, the effect of reaction (9) diminishes, influencing the observed rate less than 1% at pressures greater than 100 Torr at 298 K. The observed rate was influenced by less than 3% at all other temperatures and pressures examined in the present experiment. All reported k1 values have taken this correction into account.

1.2.2 Effect of reaction (4) on IR and UV signals

Because ^{2 o}

2 max

[NO ]

[HO ] > 60, the loss of HO2 via reaction (1) was treated as first- order. To analyze the decay of IR and UV signals, the equation

11

was used to fit the data, where s(t) was the absorbance signal (unitless) for the UV channel, and demodulated voltage signal (in V) for the IR channel. Reactions (1) and (4) are the major loss processes for HO2. Since the loss by reaction (4) was second-order, fits to the data using equation 10 resulted in values of k′ that were dependent on ^{2 o}

2 max

[NO ] [HO ] , the time span in which the fitting procedure was employed, and k1([M],T). The effect of reaction (4) on measured k1 was ascertained three different ways. First, as has been done in prior examinations⁹ of k1, kinetic modeling using FACSIMILE was used to determine the effect reaction (4) on the overall rate measurement of k1. The largest correction to k1

was a 5% decrease in the value observed at 50 Torr and 298 K. At 100 Torr, the correction was less than 3% for all temperatures. Second, measured k1 values did not differ by more than 5% when ^{2 o}

2 max

[NO ]

[HO ] was changed by an order of magnitude. Third, no significant difference in the value k1 was observed when fits were conducted over differing time intervals. This procedure was adopted because as time proceeds, the influence of reaction (4) decreases.

A positive residual baseline signal was observed in many experiments. The magnitude of this residual was typically less than 2% of the maximum HO2 signal. This residual showed negligible temporal dependence and was thus assumed to be constant for fitting purposes. The source of this residual was uncertain.

1.2.3 Effect of overlapping absorptions on UV signal

Table 1.1 lists several species and their cross sections at various UV wavelengths.

Figure 1.2 shows the contribution of these species to the UV signal at 220 nm using the FACSIMILE kinetic modeling program and the kinetic model described in Table 1.2. The model was computed at 231 K, 100 Torr total pressure of N2, [NO2] = 2 × 10¹⁵ molecules cm^-3, [CH3OH] = 3 × 10¹⁵ molecules cm^-3. H2O2 absorbance was negligible. Combined absorbances for ClONO and ClNO2 are shown. The cross section of ClONO at 220 nm is not known and was assigned the value of 2.15 × 10^-18 cm², a value measured by Molina and Molina²⁹ at 235 nm. This value was chosen because the cross section of ClONO appears to increase as wavelength decreases near 235 nm and thus would appear to be a lower limit to the actual value at 220 nm. These figures demonstrate that the acquired UV signal contains significant contributions from several species.

Despite overlapping absorptions, the HO2 rate of decay is equivalent to the first- order rate of decay of the total UV signal if equation 10 is used to fit the data and if the concentration of all species that contribute to the signal are solely dependent on reaction (1) or are constant during the time of analysis.³⁰

1.2.4 Secondary chemistry involving Cl + NO2 recombination

Possible complications arising from the formation of ClONO and ClNO2 were considered as part of the data analysis. In the present experiments, these species are formed by the reactions

13

M

2 2

Cl + NO → ClNO (12)

Buildup of these species can be significant depending on the relative concentrations of CH3OH and NO2. For example, with [CH3OH] = 2 × 10¹⁵ molecules cm^-3 and [NO2] = 4

× 10¹⁵ molecules cm^-3, the fraction of Cl that reacts with NO2 is 0.23 at 231 K and 100 Torr total pressure of N2. At 298 K, the corresponding fraction is 0.13. These calculations are based on the rate constant recommendations of DeMore et al.¹⁸ and assume that reactions (10), (11) and (12) are the only loss processes for Cl.

ClONO and ClNO2 can affect the measured rate of k1 by reacting with other species or by undergoing unimolecular processes such as decay or isomerization. Both species absorb strongly in the UV. In the present experiment, k1 was measured in the IR for [Cl]max 20% of typical values. Poor signal-to-noise prevented similar measurements in the UV. No noticeable difference between these low [Cl]max and normal [Cl]max

experiments was observed at 231 K and 298 K, indicating that the influence of Cl + NO2

on measured kinetics was negligible.

1.2.5 Sources of uncertainty

The statistical uncertainty (1σ) in the measurement of k′ due to noise in the IR signal was approximately 2%. For the UV signal, it was approximately 5%. Error in the photometric measurement of [NO2] due to changes in the flux of the NO2 gas mixture from the NO2 bulb, which was due to the chance in pressure of the bulb over the course of an experiment, was approximately 3%. Uncertainty in [CH3OH], which in turn was due to the fluctuations in measured gas flows and temperature of the bath surrounding the

methanol, was 3%. The uncertainty in the measured pressure was approximately 1% and in measured temperature was ± 1 K. The total uncertainty in the precision of the measurements of k1(T) ranged from approximately 5% to 10%. The observed fluctuations in measured k1(T) at 230 K and 219 K were observed to be somewhat higher, 8% to 15%, most likely due to the errors caused by imprecision in determining [CH3OH]. The systematic uncertainty in the cross section of NO2 at room temperature is approximately 10% at room temperature. The uncertainty in the accuracy of the vapor pressure of methanol was 5%. These systematic uncertainties are not reported in the uncertainties given for the measurements of any rate coefficients in the present work.

1.3 Results

It is demonstrated in Chapter 2 the Cl2-CH3OH-O2 source of HO2 can result in kinetic complications arising from reactions of the HO2·CH3OH complex. Studies of reaction (1) were therefore carried out to determine both the pressure and temperature dependences of k1 and to examine the possibility that complex formation enhances the observed reaction rate.

1.3.1 Effect of CH3OH on reaction (1)

The effect of CH3OH on the rate of reaction (1) was studied at 100 Torr for six different temperatures ranging from 231 K to 298 K. In addition, at 231 K, the reaction was studied at 50 Torr and 200 Torr. The observed first-order rate constant, k′, for the

15

[NO2] and [CH3OH]. Plots of [HO2] versus time at 100 Torr, 231 K, and [NO2] = 2.80 × 10¹⁵ molecules cm^-3 are shown in Figure 1.3 for different [CH3OH].

These experiments showed that there is a significant dependence on the apparent rate of reaction (1) with [CH3OH]. The dependence of k′ on [CH3OH] for various [NO2] is shown in Figure 1.4 at 100 Torr and 231 K. The dependence of k′ on [CH3OH] was well described by the equation

o [CH OH]3

k′ = ′ + ⋅k k′′ (13)

Both ko′ and k″ were both observed to be dependent on [NO2]. The value of ko′ represented the first-order rate constant at zero [CH3OH]. The value of k″ represented the dependence of the measured first-order rate constant on [CH3OH].

The trend of ko′ versus [NO2] was observed to be linear at all temperatures and was analyzed with the equation

2 2

o HO +HO 1 [NO ]2

k ′ =k + ⋅k (14)

where kHO2+HO2 represents the contribution of the HO2 + HO2 reaction to the measurement of k′o. The slope, k1, is the rate constant for HO2 + NO2 in the limit of zero methanol.

Figure 1.5 is a plot of k′o versus [NO2] at 100 Torr for 231 K, 250 K, and 298 K. The temperature dependence of k1 at 100 Torr is shown in Figure 1.6. The values of k1 are tabulated in Table 1.3. The measured values of k1 in the present study were within 3% of

the NASA values at 298 K and 288 K and were approximately 14% lower than the NASA values at 231 K, as indicated by the last column of Table 1.3.

The trend of k″ versus [NO2] was more difficult to discern. Plots of k″ versus [NO2] are shown in Figure 1.7 for 100 Torr and the temperatures, 231 K, 250 K, and 298 K. In general, the trend appeared linear and was thus described by the equation

†

o [NO ]2

k′′= + ⋅k′′ k (15)

where k^† and ko″ and represent the enhancement of k′ of enhancement of reaction (1) and reaction (4) by CH3OH respectively. The values of k^† and ko″ are listed in Table 1.3. At 288 K and 298 K, the uncertainty in the fitted value for ko″ was greater than the value itself. The temperature dependences of k^† and ko″ were analyzed with the Arrhenius equation k(T) = Ao⋅exp[(Ea/R)/T]. Plots of k^† and ko″ versus T ^-1 are shown in Figure 1.8.

For k^†, Ao = (1.6 ± 0.9) × 10^-36 cm⁶ molecule^-2 s^-1 and Ea/R = (-4360 ± 140) K. For ko″, they were (1.9 ± 3.0) × 10^-22 cm³ molecule^-1 s^-1 and (-4760 ± 370) K, respectively.

Using the above relations, k′ can be approximated at any [NO2] and [CH3OH] by the equation

2 2

†

HO +HO (4.1) [NO ]2 o [CH OH]3 [NO ] [CH OH]2 3

k′ =k +k ⋅ + ⋅k′′ + ⋅k ⋅ (16)

The effect of CH3OH on the study of the HO2 + NO2 reaction was not accounted for in previous studies. In previous studies, the rate constant of reaction (1) was equated to the

17

†

1 3

2

[CH OH]

[NO ]

dk k k

d ′ = + ⋅ (17)

A plot of

[NO ]2

dk d

′ versus temperature at [CH3OH] = 3 × 10¹⁵ molecules cm^-3 and

100 Torr is depicted in Figure 1.6. This [CH3OH] was chosen because it falls within the range used in previous experiments. As can be seen from Figure 1.6, at 231 K, 100 Torr,

[NO ]2

dk d

′ is a factor of 2.0 larger than k1 measured in the present experiment and a factor

of 1.7 larger than the NASA recommended values. The significant differences between

[NO ]2

dk d

′ from this experiment and the NASA recommended values, which were based on

experiments where [CH3OH] ranged from (2 to 8) × 10¹⁵ molecules cm^-3 suggests that measurements of k1 in which HO2 is monitored in the UV and IR differ at low temperatures.

At 231 K, the methanol enhancement of reaction (1) was investigated at 50 Torr, 100 Torr, and 200 Torr. Measurements of k1 and k^† at 100 Torr and 200 Torr for 231 K are listed in Table 1.3. Three separate attempts were made at measuring k1 and k^† at 50 Torr. For all three attempts, measured values of k′ as a function of [CH3OH] were highly scattered and not well described by equation (13). It is unclear why this was so.

The values measured at 50 Torr and 231 K are not included in our analysis.

The parameter governing the enhancement of reaction (1) by CH3OH, k^†, was increased slightly with pressure, from (2.1 ± 1.4)×10^-29 cm⁶ molecule^-2 s^-1 to (2.3 ± 1.9)

×10^-29 cm⁶ molecule^-2 s^-1 but the difference was well within the error estimates. However, the uncertainties of the measurements make it unclear whether or not there is a pressure effect on k^†. More studies over a larger range of pressures are needed. The increase in the value of k1/kJPL with pressure, noted in the last column of Table 1.4, suggests that k^† may be pressure dependent.

1.3.2 Comparison of IR and UV data

Comparisons of simultaneously acquired IR and UV signals at 100 Torr are shown in Figures 1.9.1 and 1.9.2 for 298 K and 231 K respectively. At 231 K, the IR signal indicates that HO2 is no longer present after 3 ms; however, the UV signal is non- zero and time-dependent after 3 ms. This strongly suggests that the UV channel is sensitive to species which can interfere with the HO2 absorption signal. Measurement of k′ using 231 K UV data and equation (10) is complicated by the lack of a stable baseline UV signal after all the HO2 has reacted (denoted post-HO2 signal). Despite this, the data acquired at 231 K were analyzed with equation (10) over the time span of 3 ms. Plots of k′ versus [NO2] for both the IR and UV data are shown in Figure 1.10 for 298 K and 231 K. At 298 K, IR and UV measurements agree. At 231 K, there is significant disagreement.

At 231 K, a rise (from a negative absorbance towards zero absorbance) in the 400 nm post-HO2 signal occurred simultaneously with the decrease in the 220 nm post- HO2 signal. This suggested that NO2 was generated from a temporary NO2 reservoir. The 400 nm post-HO2 signal, which has a negative absorbance value due to the consumption

19

absorbance and then became constant. The rate of increase in the 400 nm signal was similar to the rate of decrease in signal at 220 nm. These observations can be explained by the process

2 2 2 4

NO + NO ←→ N O (18)

Both N2O4, and NO2, σ^NO_220nm² = (4.7 ± 0.3) × 10^-19 cm² molecule^-1,¹⁹ contribute to the observed absorbance At 220 nm, the cross section of N2O4 is larger than the cross section for NO2 (see Table 1.2) and the time-dependent signal is dominated by the loss of N2O2

via reaction (-18). At 400 nm, NO2 absorbance dominates and the time-dependent signal is mainly due to the gain of NO2 from reaction (-18).

To illustrate the absorbance change at 400 nm, consider a typical experiment in which [NO2]eq = (1.9 ± 0.1) × 10¹⁵ molecules cm^-3. Under these conditions, [N2O4]eq = (5.8 ± 0.4) × 10¹⁴ molecules cm^-3.¹⁵ Each photolysis pulse removed (7.0 ± 0.5) × 10¹³ molecules cm^-3 of NO2, mainly due to reaction (1). In order for the system to reach equilibrium after the photolysis pulse, [NO2] increases by nearly 4 × 10¹³ molecules cm^-3 from dissociation of N2O4. The change in absorbance in the 400 nm post-HO2 signal is about 0.003 absorbance units, which is measurable in the present experiment.

The value of k-18 was measured to be (36 ± 10) s^-1. This compares favorably with previous measurements³¹ of k-18 made at higher temperatures which predict values of k-18

between 20 s^-1 and 180 s^-1 at 231 K and 100 Torr.

Attempts were made to study reaction (1) in the absence of hydrogen-bonding species. A gas mixture of F2/H2/O2/N2 was flowed into the cell at concentrations (molecules cm^-3) of F2: 5 × 10¹⁶; H2: 5 × 10¹⁶; O2: 5 × 10¹⁷ and balance N2 at 100 Torr and 231 K. When NO2 was added to the gas mixture, an unexpected explosion took place.

This may have occurred as a result of a thermal, wall-catalyzed, reaction or from photolysis by ambient light. The reaction mixed was judged to be sufficiently unstable that no further studies were conducted using the F2-H2 system.

1.3.3 Measurements of k1 at low [CH₃OH]

A second set of measurements of the rate of reaction (1) was obtained with [CH3OH] = 2.0 × 10¹⁴ molecules cm^-3. These measurements were done between 50 Torr and 200 Torr, and between 219 K and 295 K. As stated above, complications from Cl + NO2 were found to be insignificant. Results from the first set of experiments indicated that at 231 K, 200 Torr, and [CH3OH] = 2.0 × 10¹⁴ molecules cm^-3, the calculated value of

[NO ]2

dk d

′ was approximately 5% greater than the measured value of k1. At 219 K, the

observed rates were calculated to be nearly 15% higher than k1. The observed values from this second set of measurements, corrected for the presence of CH3OH, are shown in Figure 1.11.

1.4 Discussion

1.4.1 Quantifying the results

21

2 -1 o( ) [M]

1+ log o ( )

1

o

( ) [M]

([M], )

( ) [M]

1 ( )

k T k T c

k T k T F

k T k T

∞

   ⋅  

    

 

  

   

 

∞

= + ⋅ ⋅ ⋅ (19)

has been adopted by the NASA and IUPAC data evaluation panels to describe the falloff behavior of association and unimolecular decomposition reactions. The parameters ko(T) and k_∞(T) are the low and high pressure limiting rate constants, respectively with their temperature dependences given by

ko(T) = ko(300K)⋅(T/300)^-m (20)

and

k_∞(T) = k_∞(300K)⋅(T/300)^-n (21)

The parameter Fc was assigned the value 0.6 in accordance with the procedure adopted by the NASA data evaluation panel (ref). The parameters that were acquired in the fitting process were ko(300K), k_∞(300K), m, and n. Two fitting trials are tabulated in Table 1.4.

Trial 1 employed both sets of data from the present experiment. Trial 2 employed all the data of Trial 1 and also that acquired at T ≥ 277 in experiments by Sander and Peterson⁸ and Kurylo and Ouellette.^9,10 Both trials were weighted by the stated uncertainties. As Table 1.4 shows, for each fitted value, the discrepancy between Trial 1 and Trial 2 is greater than the combined uncertainty.

The lack of agreement between the trials may result from insufficient parameterization of equation (19). Since one of the principal aims of this paper is to provide a description of reaction (1) that is useful for atmospheric modeling, and since the adopted procedure by the NASA data evaluation panel has adopted the use of the equation (19), further parameterization was not adopted.

As stated in the introduction, impact of reaction (1) on atmospheric chemistry is greatest from the upper troposphere to the middle stratosphere. An assessment of how each of the trials describes reaction (1) in this region of the atmosphere can be quantified by comparing the calculated rate from each of the trials to the current NASA recommended value at 231 K and 100 Torr. This has been done in Table 1.5. Both Trials calculate rates that are about 10% lower than the current recommendation. The measured value from data set 1 was 10% lower than currently recommended. Data set 1 is highlighted because was a direct measurement of k1 at 231 K and 100 Torr.

The above analysis indicates that for modeling the chemistry of the upper troposphere to the middle stratosphere, there is little difference between Trial 1 and Trial 2 though the parameters acquired from Trial 2 best describe reaction (1) over the widest range of pressures and temperatures. Figure 1.12 depicts the difference between Trial 2 and the current recommended values.

1.4.2 Enhancement by CH3OH

The observed enhancement of reaction (1) by CH3OH can be explained by the reaction sequence

23

M

2 2 2 2

HO + NO → HO NO (1)

3 2 3 2

CH OH + HO ←→ CH OH HO⋅ (22)

3 2 2

CH OH HO + NO ⋅ → Products (23)

If the steady-state approximation is used for [CH3OH⋅HO2],

[NO ]2

dk d

′ = k1 + 2⋅k23⋅K22⋅[CH3OH] (24)

where

[NO ]2

dk d

′ is the observed rate constant discussed above and K22 describes the

equilibrium between HO2, CH3OH, and CH3OH⋅HO2. From the present experiment, the rate enhancement of reaction (1), k^†,was measured to be (1.6 ± 0.9) × 10^-36 × exp((4360 ± 140)/T) cm⁶ molecule^-2 s^-1. In a prior study, the enhancement of reaction (1) by CH3OH was described in a similar fashion and measured to be (2.5 ± 2.4) × 10^-36 × exp(-4570 ± 120) cm⁶ molecule^-2 s^-1.¹⁷ If it is assumed that the rate of reaction (23) depends very little on temperature, then the temperature dependence of the enhancement can be shown result from the enthalpy change of Keq. In the present study, the enthalpy change for reaction (1) was measured to be -(8.66 ± 0.28) kcal mol^-1. For reaction (4), the enthalpy change was measured to be -(9.1 ± 0.2) kcal mol^-1. Both these values correspond to strong hydrogen bonded complexes.

The similarity in enhancement between reactions (1) and (4) suggests that the process HO2 + CH3OH⋅X → Products, where X = HO2 or NO2, may occur at near collision frequency as has been suggested by prior researchers.¹⁶

1.4.5 Conclusion

The effect of methanol on the observed rate of HO2 + NO2 was measured. This information was used to measure the rate constant of HO2 + NO2 in limit of zero methanol k1. IR spectroscopy was employed, minimizing the influence of the equilibrium between NO2 and N2O4 in determining the rate, a process not taken into account in prior studies. The results indicated that at temperatures lower than 250 K, k1 was lower than the current NASA recommended values. At 231 K, 100 Torr, k1 was nearly 10% lower.

Parameterizations of the rate of k1 using a simplified Troe termolecular equation was done using the present data in addition to that taken by prior researchers. Only data that in which the effect of CH3OH was minimal was included. It was found that the simplified equation did not adequately describe all the data. However, it did describe the rate of reaction (1) in the pressure and temperature regime of importance to atmospheric chemistry.

The methanol effect was analyzed and found to be remarkable similar to that for the enhancement of the HO2 + HO2 system. This suggests that current models discussed in the literature approximate the process well.

25

Acknowledgements

This research was carried out by the Jet Propulsion Laboratory, California Institute of Technology, under contract with the National Aeronautics and Space Administration.

Support is acknowledged from the NASA Upper Atmosphere Research and Tropospheric Chemistry Programs. This research has also been supported in part by a grant from the U.S. Environmental Protection Agency National Center for Environmental Research’s Science to Achieve Results (STAR) program, through grant R826236-01-0. It has not been subjected to any EPA review and therefore does not necessarily reflect the views of the Agency, and no official endorsement should be inferred. We would like to acknowledge the technical support of Dave Natzic, Jürgen Linke, Siamak Forouhar, Dave Dougherty, and Sam Keo of JPL.

1.5 References

1. Simonaitis, R. and J. Heicklen J. Phys. Chem 78: 653 (1974).

2. Cox, R. A. and R. G. Derwent J. Photochem. 4: 139 (1975).

3. Simonaitis, R. and J. Heicklen J. Phys. Chem 80: 1 (1976).

4. Howard, C. J. "Kinetics of the reaction of HO2 with NO2." J. Chem. Phys. 67: 5258 (1977).

5. Niki, H., P. Maker, et al. "FTIR of PNA from HO2 + NO2." Chemical Physics Letters 45: 564 (1977).

6. Simonaitis, R. and J. Heicklen Int. J. Chem. Kinet. 10: 67-87 (1978).

7. Cox, R. A. and R. Patrick Int. J. Chem. Kinet. 11: 635 (1979).

8. Sander, S. P. and M. Peterson "HO2 + NO2." J. Phys. Chem. 88: 1566-1571 (1984).

9. Kurylo, M. J. and P. A. Ouellette "HO2 + NO2." J. Phys. Chem. 90: 441-444 (1986).

10. Kurylo, M. J. and P. A. Ouellette "Rate Constants for the Reaction HO2 + NO2 + N2 −

> HO2NO2 + N2: The Temperature Dependence of the Falloff Parameters." J.

Phys. Chem. 91: 3365-3368 (1987).

11. WMO (1983). The Statosphere: 1981, NASA.

12. Rinsland, C. P., R. Zander, et al. "Evidence for the Presence of the 802.7 cm^-1 Band Q Branch of HO2NO2 in High Resolution Solar Absorption Spectra of the Stratosphere." Geophysical Research Letters 13: 761-764 (1986).

13. Sen, B., G. C. Toon, et al. "Measurements of Reactive Nitrogen in the Stratosphere."

Journal of Geophysical Research-Atmospheres 103: 3571-3585 (1998).

14. Atkinson, R., D. L. Baulch, et al. "Summary of Evaluated Kinetic and Photochemical Data for Atmospheric Chemistry - Web Version December 2000." (2000).

15. Sander, S. P., R. R. Friedl, et al. (2000). Chemical Kinetics and Photochemical Data for Use in Stratospheric Modeling, Evaluation Number 13. Pasadena, CA, Jet Propulsion Laboratory, California Institute of Technology.

16. Andersson, B. Y., R. A. Cox, et al. "The Effect of Methanol on the Self Reaction of HO2 Radicals." Int. J. Chem. Kinetics 20: 283-295 (1988).

17. Christensen, L. E., S. P. Sander, et al. "Kinetics of HO2 + HO2 → H2O2 + O2: Implications for Stratospheric H2O2." Geophysical Research Letters (2002).

18. DeMore, W. B., S. P. Sander, et al. (1997). Chemical Kinetics and Photochemical Data for Use in Stratospheric Modeling, Evaluation Number 12. Pasadena, CA, Jet Propulsion Laboratory, California Institute of Technology.

27

20. Davidson, J. A., C. A. Cantrell, et al. J. Geophys. Res. 93: 7105-7112 (1988).

21. Herriott, D. and H. Schulte "Folded Optical Delay Lines." Appl. Optics 4: 883-889 (1965).

22. Trutna, W. and R. Byer "Multiple-pass Raman gain cell." Appl. Optics 19: 301-312 (1980).

23. Monsour, J. (2001). Private communication.

24. Tuckett, R. P., P. A. Freedman, et al. "The emission bands of HO2 between 1.43 and 1.51 microns." Molecular Physics 37: 379-401 (1979).

25. Hunziker, H. E. and H. R. Wendt J. Chem. Phys. 60: 4622 (1974).

26. Johnson, T. J., F. G. Wienhold, et al. J. Phys. Chem 95: 6499-6502 (1991).

27. Kircher, C. C. and S. P. Sander "Kinetics and Mechanism of HO2 and DO2

Disproportionations." J. Phys. Chem. 88: 2082-91 (1984).

28. Curtis, A. R. and W. P. Sweetenham (1987). FACSIMILE/CHEKMAT, H015 ed.

Harwell: Oxfordshire (UK).

29. Molina, L. T. and M. J. Molina Geophys. Res. Lett. 4: 83-86 (1977).

30. Sander, S. P. and R. T. Watson J. Phys. Chem. 84: 1664 (1980).

31. Markwalder, B., P. Gozel, et al. J. Chem. Phys. 97: 5472-5479 (1992).

Table 1.1. Cross sections for various species.

species σ220nm ^a σ225nm ^a σ230nm ^a σ400nm ^a ref.

HO2 3.41 2.88 2.30 18

NO2 0.47 0.39 0.28 0.60 19

N2O4 6.68 4.11 2.55 19

HO2NO2 1.18 0.94 0.79 18

H2O2 0.26 0.22 0.18 18

Cl2 0.02 18

ClNO2 3.39 2.83 2.26 29

a units are 10^-18 cm²

29

Table 1.2. Relevant reactions.

Reaction Reference Cl + CH3OH → HCl + CH2OH NASA

CH2OH + O2 → HO2 + CH2O NASA HO2 + HO2 → H2O2 + O2 My GRL

NO + HO2 → NO2 + OH NASA

OH + NO2 → HNO3 NASA OH + HO2 → H2O + O2 NASA

Cl + HO2 → HCl + O2 NASA HO2 + NO2 + M → HO2NO2 + M NASA

Cl + NO2 + M → ClONO + M NASA Cl + NO2 + M → ClNO2 + M NASA

Cl + ClONO → Cl2 + NO2 NASA Cl + ClNO2 → Cl2 + NO2 NASA NO2 + NO2 + M → N2O4 + M This expt.

N2O4 + M → NO2 + NO2 + M This expt.

Table 1.3. Fitted values at different temperatures.

T (K) P ^a k1 b

k^†c ko″ ^d k1/kJPL

298 100 4.0±0.1 0.24±0.06 too noisy 1.00 288 100 4.4±0.1 0.58±0.09 0.9±2.0 0.98 273 100 5.1±0.1 1.5±0.2 4.5±3.2 0.93 250 100 6.7±0.1 6.3±0.2 41±2 0.92 240 100 6.9±0.1 13.8±0.4 80±5 0.83 231 100 8.5±0.3 20.7±1.4 160±10 0.90 231 200 9.4±0.4 23.0±1.9 200±60 0.66

a units are Torr

b units are 10^-13 cm³ molecule^-1 s^-1

c units are 10^-29 cm⁶ molecule^-2 s^-1 d units are 10^-15 cm³ molecule^-1 s^-1

31

Table 1.4. Fitted values for Troe equation.

Trial koa n k_∞^b m fit/kNASA

1 2.4 ± 0.1 2.1 ± 0.3 1.9 ± 0.1 4.2 ± 0.4 0.93 2 1.9 ± 0.1 3.7 ± 0.2 2.9 ± 0.1 1.1 ± 0.3 0.89

a units are 10^-31 cm⁶ molecule^-2 s^-1

b units are 10^-12 cm³ molecule^-1 s^-1

Figure 1.1. Experimental apparatus.

33

0.03

0.02

0.01

0.00

Change in Absorbance (unitless)

20 15

10 5

0

Time (ms) HO₂

HO₂NO₂

NO₂ N₂O₄

Total

ClNO₂ + ClONO

Figure 1.2. Simulated UV absorbances at 369.50 nm using FACSIMILE.

1.6

1.4

1.2

1.0

0.8

0.6

0.4

0.2

0.0 [HO2] (1013 molecules cm-3 )

4 3

2 1

0 -1

time (ms)

[CH₃OH]=3.84E15 [CH₃OH]=7.25E14

Figure 1.3. Decay of [HO2] due to the HO2 + NO2 reaction at different [CH3OH] at 231 K, 100 Torr.

35

7000

6000

5000

4000

3000

2000

1000

0 k' (s-1 )

5x10¹⁵ 4

3 2

1 0

[CH₃OH] (molecules cm^-3) [NO₂] = 3.71E15

[NO₂] = 1.70E15

[NO₂] = 5.78E14

Figure 1.4. k′ versus [CH3OH] for various [NO2] at 231 K, 100 Torr.

3500

3000

2500

2000

1500

1000